Chapter 6 The Periodic Table and

Periodic LawThe elements, which make up all

living and non-living matter, fit into a orderly table. When interpreted

properly, the table describes much of the elements physical and chemical

properties.

What is the Periodic Law and how was it formulated?

• Demitri Mendeleev is known as the father of the periodic table

• He arranged the elements in families (groups) and periods (rows,series) according to atomic mass and properties

• Mendeleev noted that the chemical and physical values for elemental properties would either be high or low depending upon the group under observation.

• He proposed the first Periodic Law "The properties of the elements are a periodic function of their atomic masses"

• Left blanks in his table for undiscovered elements

Moseley’s Modern-Periodic Law

• There was some inconsistencies with Mendeleev’s table

• In the early 1900’s Moseley was able to experimentally determine the atomic number of all known elements

• Moseley then proceeded to rearrange the elements according to increasing atomic numbers.

• New/Modern Periodic law states that the properties of elements are a periodic function of their atomic number

The Modern Periodic Table

• Glenn T. Seaborg won the Nobel Prize for his work in nuclear chemistry

• In 1944, formulated the “actinide concept” of heavy element electronic structure. This concept predicted that the fourteen actinides

Some Characteristics of Groups

Group 1 (IA) - Alkali Metals

( metal characteristics - shiny, malleable, ductile, good conductors)

• Very active metals - activity increases as you go down group

• All have one valence electron - form +1 cations by losing an electron

• react violently with water

Group 2 (IIA) - Alkaline Earth Metals

• activity increases as you move down the column not as reactive as alkali metals

• Ca, Sr, and Ba react violently when they come into contact with water

• All have two valence electrons

• Form +2 cations by losing 2 electrons

Group 17 (VIIA) - Halogens

• All gain one electron to form anions with a charge of -1.

• All are nonmetals except for At which is a semimetal

• All are diatomic in their elemental form

Group 18 (VIIIA) - Noble (Rare) Gases

• Mistakenly labeled as "inert gases" until about 30 years ago because it was thought that these gases did not react with anything.

• Noble gases have filled valence (outermost) shells.

Periodic Trends

• Atomic Radii

1) As you move down a group, atomic radius increases.

2) WHY? - The number of energy levels increases as you move down a group as the number of electrons increases.

• As you move across a period, atomic radius decreases.

• WHY? - As you go across a period, electrons are added to the same energy level. At the same time, protons are being added to the nucleus. The concentration of more protons in the nucleus creates a "higher effective nuclear charge."

First Ionization Energy

• Definition: The energy required to remove the outermost (highest energy) electron from a neutral atom in its ground state.

1) As you move down a group, first ionization energy decreases.

2) WHY? Electrons are further from the nucleus and thus easier to remove the outermost one - + shielding effect of other electrons

3) As you move across a period, first ionization energy increases.

4) WHY? - As you move across a period, the atomic radius decreases, the attraction from the positive nucleus gets larger

Electronegativity

• Definition: The ability of an element to attract electrons in a chemical bond

1) As you move down a group,

electronegativaty decreases

2) As you move across a period, it

increases.